2. Chapter 8: Covalent Compounds

3. Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

4. Polarity of Molecules Dipole: def: two charges, equal in magnitude and opposite in sign, are separated by a distance Polarity of Polyatomic Molecules Each bond can be polar. The orientation of these polar bonds determines whether the molecule is polar overall. It is possible for a molecule with polar bonds to be either polar or non-polar.

5. Polarity of Molecules Dipole Moments of Polyatomic Molecules Example: in CO 2, each C-O dipole is canceled because the molecule is linear. In H 2 O, the H-O dipoles do not cancel because the molecule is bent.

6. Polarity of Molecules Dipole Moments of Polyatomic Molecules

7. Models  Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.  Properties of Models:  A model does not equal reality.  Models are oversimplifications, and are therefore often wrong.  Models become more complicated as they age.  We must understand the underlying assumptions in a model so that we don’t misuse it.

8. Lewis Structure  Shows how valence electrons are arranged among atoms in a molecule.  Valence electrons are found in the outermost energy level of an atom. They are involved in bonding.  Reflects the central idea that the stability of a compound relates to noble gas electron configuration.

9. Lewis Dot Structure Rules:  Treat ions separately (e.g. NH 4 Cl)  Count only valence electrons  Assemble bonding framework  Fill up non-bonding electrons on outer atoms  Fill up non-bonding electrons on inner atoms  Calculate Formal Charge  Minimize Formal Charge

10. To Complete a Lewis Structure:  Must be able to recognize polyatomic ions  Must be able to identify valence electrons  Must be able to construct Bond framework

11. Hints on Lewis Dot Structures 1. Octet “rule” is the most useful guideline. 2. Carbon always forms 4 bonds. 3. Other 2nd row elements (N, O, F) observe the octet rule. 4. Hydrogen forms one bond with other atoms to complete its octet with 2 electrons. 5. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 6. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.

12. Notes About the Octet Rule 7. When multiple bonds are forming, they are usually between C, N, O or S. 8. Nonmetals can form single, double, and triple bonds, but not quadruple bonds. 9. Always account for single bonds and lone pairs before forming multiple bonds. 10. Look for resonance structures. 11. When writing Lewis structures, satisfy octets first, then place electrons around central atoms having available d orbitals.

13. PCl 3 Bonding Pairs Lone Pairs (a.k.a. nonbonding electrons) 5+(3*7)=26 e -

14. Try Some Examples:  C 2 H 5 OH  CH 3 CH 2 NH 2  Cl 2 CO  Ozone (O 3 )  NO 2

15. Formal Charge Difference between the # of valence electrons in the free atom and the # of electrons assigned to that atom in the Lewis structure. FC = formal charge; G.N. = Group Number #BE = bonding electrons; #LPE = lone pair electrons If Step 4 leads to a positive formal charge on an inner atom beyond the second row, shift electrons to make double or triple bonds to minimize formal charge, even if this gives an inner atom with more than an octet of electrons.

16. Formal Charge Not as good Better

17. Multiple Bonds It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): One shared pair of electrons = single bond (e.g. H 2 ); Two shared pairs of electrons = double bond (e.g. O 2 ); Three shared pairs of electrons = triple bond (e.g. N 2 ). Generally, bond distances shorten with multiple bonding. Covalent Bonding Octet in each case

19. Resonance  Occurs when more than one valid Lewis structure can be written for a particular molecule.

20. Odd Number of Electrons… NO Number of valence electrons = 11 NO 2 Number of valence electrons = 17 Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons) Molecules and atoms which are neutral (contain no formal charge) and with an unpaired electron are called Radicals

21. Beyond the Octet  Elements in the 3 rd period or higher can have more than an octet if needed.  Atoms of these elements have valence d orbitals, which allow them to accommodate more than eight electrons.

22. More than an Octet… PCl 5 Elements from the 3rd period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding P : (Ne) 3s 2 3p 3 3d 0 Number of valence electrons = 5 + (5 x 7) = 40 SF 4 S : (Ne) 3s 2 3p 4 3d 0 Number of valence electrons = 6 + (4 x 7) = 34 The Larger the central atom, the more atoms you can bond to it – usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency.

23. Less than an Octet… BCl 3 Group 13 atom only has six electrons around it

24. Molecular Shapes Lewis structures give atomic connectivity: they tell us which atoms are physically connected together. They do not tell us the shape. The shape of a molecule is determined by its bond angles. Consider CCl 4 : experimentally we find all Cl- C-Cl bond angles are 109.5 . Therefore, the molecule cannot be planar. All Cl atoms are located at the vertices of a tetrahedron with the C at its center.

25. Molecular Shape of CCl 4

26. VSEPR Theory  In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts whichever 3D geometry minimized this repulsion.  We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.

27. Why is VSEPR Theory Important?  Gives a specific shape due to the number of bonded and non- bonded electron pairs in a molecule  Tells us the actual 3-D structure of a molecule  In bonding, electron pairs want to be as far away from each other as possible.

28. VSEPR and Resulting Geometries

29. How does VSEPR THEORY work? We can use VSEPR theory using 4 steps. 1. Draw the Lewis Structure for the molecule. Example: SiF 4 F F-Si-F F

30. How does VSEPR THEORY work? We can use VSEPR theory using 4 steps 1. Draw the Lewis Structure for the molecule. 2. Tally the number of bonding pairs and lone (non-bonding) pairs on the center atom. F F-Si-F F Bonding pairs: 4 Lone pairs on central atom: 0

31. How does VSEPR THEORY work? We can use VSEPR theory using 4 steps 1. Draw the Lewis Structure for the molecule 2. Tally the number of bonding pairs and lone pairs on the center atom. 3. Arrange the rest of the atoms so that they are as far away from each other as possible. Si F F F F

32. How does VSEPR THEORY work? We can use VSEPR theory using 4 steps 1. Draw the Lewis Structure for the molecule 2. Tally the number of bonding pairs and lone pairs on the center atom. 3. Arrange the rest of the atoms so that they are as far away from each other as possible 4. Give the type of geometry the molecule has: Tetrahedral

33. Another Example: To determine the electron pair geometry: 1) draw the Lewis structure; 2) count the total number of electron pairs around the central atom. 3) arrange the electron pairs in one of the geometries to minimize e  -e  repulsion. 4) multiple bonds count as one bonding pair for VSEPR

34. The VSEPR Model Predicting Molecular Geometries

35. The VSEPR Model Predicting Molecular Geometries

36. The VSEPR Model Difference between geometry and shape Geometry: We determine the geometry only looking at electrons. All the atoms that obey the octet rule have the same tetrahedral-like geometry. Shape: We name the shape by the positions of atoms. We ignore lone pairs in the shape.

37. The VSEPR Model Predicting Shape Shape

38. The VSEPR Model Predicting Shape Shape

39. The VSEPR Model The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles By experiment, the H-X-H bond angle decreases on moving from C to N to O: Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs. Therefore, the bond angle decreases as the number of lone pairs increase.

40. The VSEPR Model The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles Similarly, electrons in multiple bonds repel more than electrons in single bonds.

41. The VSEPR Model Molecules with Expanded Valence Shells Atoms that have expanded octets have AB 5 (trigonal bipyramidal) or AB 6 (octahedral) electron pair geometries. Examples: PF 5 trigonal bipyramidalSCl 6 octahedral

42. The VSEPR Model Molecules with Expanded Valence Shells

43. The VSEPR Model Molecules with Expanded Valence Shells

44. The VSEPR Model Molecules with More than One Central Atom In acetic acid, CH 3 COOH, there are three central atoms. We assign the geometry about each central atom separately.

45. Hybrid Orbitals  In bonding, s and p orbitals are used in bonding. It is easy to tell which ones are used by looking at our molecule.  For example, CH 4. Looking again at the Lewis structure, we see that there are 4 bonds. We call this sp 3 hybridized.

46. Hybrid Orbitals  Regions of electron density-EACH BOND AND LONE PAIR OF ELECTRONS ON THE CENTRAL ATOM IS KNOWN AS A REGION OF ELECTRON DENSITY.  2 regions of electron density-sp hybridized  3 regions of electron density-sp 2 hybridized  4 regions of electron density-sp 3 hybridized

47. Hybridization sp Hybrid Orbitals The two lobes of an sp hybrid orbital are 180  apart.

48. Hybrid Orbitals sp 2 Hybrid Orbitals Important: when we mix n atomic orbitals we must get n hybrid orbitals. sp 2 hybrid orbitals are formed with one s and two p orbitals. (Therefore, there is one unhybridized p orbital remaining.) The large lobes of sp 2 hybrids lie in a trigonal plane. All molecules with trigonal planar electron pair geometries have sp 2 orbitals on the central atom.

49. Hybridization

50. sp 3 Hybrid Orbitals sp 3 Hybrid orbitals are formed from one s and three p orbitals. Therefore, there are four large lobes. Each lobe points towards the vertex of a tetrahedron. The angle between the large lobes is 109.5  All molecules with tetrahedral electron pair geometries are sp 3 hybridized.

51. Hybridization

52. Hybrid Orbitals

54. Summary To assign hybridization: 1.Draw a Lewis structure. 2.Assign the geometry using VSEPR theory. 3.Use the geometry to determine the hybridization. 4.Name the shape by the positions of the atoms.

55. Hybridization and Multiple Bonds Multiple bonds overlap differently and are called  -bonds and  -bonds All single bonds are  Double bonds contain 1  and 1  bond Triple bonds contain 1  and 2  bonds

56. Bond Energy

57. Covalent Bonding & Orbital Overlap As two nuclei approach each other their atomic orbitals overlap. As the amount of overlap increases, the energy of the interaction decreases. At some distance the minimum energy is reached. The minimum energy corresponds to the bonding distance (or bond length).

58. Covalent Bonding & Orbital Overlap As the two atoms get closer, their nuclei begin to repel and the energy increases. At the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).

59. Bond Energies  Bond breaking requires energy (endothermic).  Bond formation releases energy (exothermic). H = D( bonds broken )  D( bonds formed ) energy required energy released

60. Bond Energies